8.6,8.7 Periodic Properties of the Elements

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1 Pre -AP Chemistry 8.6,8.7 Periodic Properties of the Elements READ p , Practice Problems Pg 315 -Exercise 8.9 Pg #36, 55, 64, 66, 67, 69, 72, 80 Periodic Trends are predictable patterns of physical and chemical properties of atoms based on their location on the periodic table. Remember the Periodic Law (law of chemical periodicity) states that Physical and chemical properties are periodic functions of their atomic numbers We will cover the following periodic trends: 1. Atom Size 2. Ion Size 3. Ionization Energy 4. Electron Affinity 5. Metallic Character 6. Electronegativity 7. Effective nuclear charge* 8. Shielding effect * (not in text directly) Size of Atoms & Ions / Effective Nuclear Charge Atom Size It's hard to determine where the "edge" of an atom is since the "edge" is part of the electron cloud (which is only a probable location of electrons) It's easier to look at diatoms (molecules composed of 2 of the same kind of atoms bonded together like H 2, O 2, or Cl 2) and measure the radius of the atom as 1/2 the distance between the two bonded nuclei. This is called the "bonding atomic radius" and it is slightly smaller than a non-bonding atom's radius. Atomic Radii have a very specific trend on the periodic table as shown in the picture below: DECREASE TREND: Atomic Radius decreases as you move up and to the right on the periodic table. One way to explain this trend is to look at the effective nuclear charge of atoms Shielding Effect An outer electron might be "shielded" from the attractive force of protons if there are other electrons

2 between it and the proton TREND: Shielding effect increases as you move down a group with the addition of more energy levels and increased inner electrons shielding the outer electrons from attractive force of protons Shielding effect stays the same across a period as more electrons are added to same energy level Effective Nuclear Charge (pgs ) This is the force of attraction felt between a proton and an outer electron Effective Nuclear Charge = (Actual Nuclear Charge) - (Shielding Effect) To Calculate the Effective Nuclear Charge (Z eff) : Z eff = Z - S where Z is the # of protons and S is the # of core electrons *As Z eff increases, there is more attraction between the nucleus and an outer electron, so the atomic radius will decrease. Example: Compare Z eff for Silicon and Phosphorus Si P Z eff = = 4 Z eff = = 5 Since Z eff is larger for P, P will have a smaller radius (there is more attractive force between the protons and outer electrons in phosphorus than there is in silicon) Example: Compare Z eff for Fe, Co, Ni (all transition metals) *It's hard to see what the core electrons are for transition metals, so it might help to do electron configurations. Remember: core electrons are all the electrons EXCEPT the valence ones. Fe Co Ni 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 Z eff=26-24=2 Z eff=27-25=2 Z eff=28-26=2 *Zeff doesn't change within a period for the transition metals so we don't see them get very much smaller as you move RIGHT within a period. QUESTIONS YOU SHOULD KNOW HOW TO ANSWER: WHY do atoms get LARGER as you move DOWN within the same group on the periodic table? As you move down a group, electrons are added to higher energy levels that are consequently farther from the nucleus. In addition, the electrons in higher energy levels are more shielded by the inner electrons thus greatly reducing the attractive force between the nucleus and outer electrons and allowing the outer electrons to move farther from the nucleus thus resulting in a larger atom. WHY do atoms get SMALLER as you move RIGHT within the same period on the periodic table? As you move right within a period, each atom has additional protons and electrons, but the additional electrons are not being added to higher principal energy levels so the amount of shielding by the inner electrons does not increase. Therefore, the additional protons in the nucleus are able to pull more strongly on the outer electrons and thus pull them in closer to the nucleus making the atom smaller. PROBLEMS YOU SHOULD KNOW HOW TO DO: Arrange each group of atoms from smallest to largest: a. Mg, S, Si b. As, N, P Ion Size This trend doesn't have much of a pattern on the periodic table, but it is definitely predictable.

3 TREND: CATIONS ARE SMALLER than the atom from which they form To form a cation, an atom must lose it's outer electrons thus getting rid of an entire principal energy level You can prove that cations are smaller based on Z eff also: Mg Mg 2+ 1s 2 2s 2 2p 6 3s 2 1s 2 2s 2 2p 6 Z eff = = 2 Z eff = 12-2 = 10 Z eff is much larger for the Mg2+ ion so that means the protons have more attraction for the outer electrons and thus make the ion smaller. TREND: ANIONS ARE LARGER than the atom from which they form To form an anion, an atom will gain electrons The additional electrons cause electron/electron repulsion which tends to "push" the electrons away from each other (they spread out more) thus making the ion larger This effect is not due to changes in Z eff since the additional electrons are added to the same principal energy level. F F - 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6 Z eff = 9-2 = 7 Z eff = 9-2 = 7 Zeff is the same for both atom and anion, so the change in size must be due to increased electron repulsions from the gained electrons. Isoelectronic Species These are particles (atoms and ions) that have the same number of electrons Example: Which is smallest? Ne Na + Mg 2+ 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 1s 2 2s 2 2p 6 10e 10e 10e Z eff=10-2=8 Z eff=11-2=9 Z eff=12-2=10 *Mg2+ is the smallest because it has the largest Z eff *For isoelectronic species, the one with the most protons is always going to be the smallest!!! PROBLEMS YOU SHOULD KNOW HOW TO DO: Arrange from smallest to largest: Fe 3+ K + S 2- Se 2-

4 Ionization Energy The minimum amount of energy required to remove an electron from a gaseous atom in its ground state 1st ionization energy = energy required to remove 1st electron 2nd ionization energy = energy required to remove 2nd electron 3rd ionization energy = energy required to remove 3rd electron Example: Na atom 1st I.E. = 496 kj/mol (results in Na+) 2nd I.E. = 4,560 kj/mol (results in Na2+) *We never see an Na2+ ion because the energy requirement to make it is too high!!! Example: Mg atom 1st I.E. = 738 kj/mol (Mg+) 2nd I.E. = 1450 kj/mol (Mg2+) 3rd I.E. = 7730 kj/mol (Mg3+) *Note how big the 3rd I.E. is...that's why you never see an Mg3+ ion. The general trend for Ionization Energy is shown below: TREND: Ionization energy increases as you move up and to the right on the periodic table. You can see that smaller atoms have higher ionization energies because the outer electrons are more strongly attracted to the nucleus. There are several exceptions to the ionization energy periodic trend: Ionization energy drops going from "s" orbitals to "p" orbitals in the same period. This is because electrons in "p" orbitals are higher energy and are thus easier to remove from the atom There is another dip in energy when you go from group 15 to group 16 in the same period. This is due to the pairing of electrons in the "p" orbitals Group 15 Group 16 2 electrons in the same orbital causes slightly more repulsion so it makes the electron easier to remove

5 Electron Affinities The energy change associated with adding an electron to an atom It is easy to add an electron to an atom that has a high electron affinity It is difficult to add an electron to an atom that has a low electron affinity Some textbooks show electron affinity values as negative energy values because electrons must release energy to get closer and closer to the nucleus and eventually become part of the atom (energy releases are negative, right?) In these books, chlorine has the most negative electron affinity value and thus has the highest electron affinity of any atom (it is the easiest atom to add an electron to) There is a more recent trend in textbooks to show electron affinity values as positive energy values which makes them a little easier to work with In these books, chlorine has the most positive electron affinity value, yet it still has the highest electron affinity of any atom (it is still the easiest atom to add an electron to) Just know that atoms with large electron affinity values (whether they are positive or negative number) will generally accept electrons. NOTE: Noble gases are not involved in this trend (because they do NOT want another electron to be added to them) TREND: Electron Affinity increases as you move up and to the right on the periodic table.

6 Additional Trends (Metallic Character & Electonegativity) Metallic Character The easier it is to remove an electron, the more "metallic" an atom is. Therefore, larger atoms have the most metallic character Because their outer electrons are so far from the nucleus, they are not strongly attracted to the nucleus TREND: Metallic character increases as you move down and left on the periodic table. Electronegativity The attraction an atom has for electrons in a bond Example: The HF molecule H and F share electrons to form a bond between them. However, Fluorine has a stronger attraction for the shared electrons and thus pulls the electrons closer to it. This creates a region of higher electron density around the fluorine nucleus and makes the "fluorine end" of this compound more negatively charged. We call this a polar molecule when there is an uneven sharing of electrons in a bond. Fluorine is the most electronegative element. Each element is assigned an electronegativity number from 0 to 4 (4 being the most electronegative) TREND: Electronegativity increases as you move up and to the right on the periodic table.